Part 1: Ionic Bonds

Welcome to Types of Chemical Bonds!

Chemical bonds hold atoms together to form compounds. The type of bond that forms depends on the properties of the atoms involved — particularly their electronegativity and whether they are metals or nonmetals.

In this seven-part series, you'll master ionic, covalent, and metallic bonding — essential knowledge for the AP Chemistry exam.

In Part 1, we focus on ionic bonds: how they form, what holds them together, and the properties they produce.

Ionic Bond Formation: Electron Transfer

An ionic bond forms when one or more electrons are transferred from a metal atom to a nonmetal atom.

The Process

1. A metal atom loses one or more valence electrons → forms a cation ($+$ charge)
2. A nonmetal atom gains those electrons → forms an anion ($-$ charge)
3. The oppositely charged ions attract each other via electrostatic (Coulombic) forces

Example: Sodium Chloride (NaCl)

$\text{Na} \rightarrow \text{Na}^+ + e^-$ $\text{Cl} + e^- \rightarrow \text{Cl}^-$

• Sodium ($Z = 11$): $[\text{Ne}]\,3s^1$ → loses 1 electron → $\text{Na}^+$ with noble gas configuration $[\text{Ne}]$
• Chlorine ($Z = 17$): $[\text{Ne}]\,3s^2\,3p^5$ → gains 1 electron → $\text{Cl}^-$ with noble gas configuration $[\text{Ar}]$

Why Does This Happen?

Metals have low ionization energies (easy to remove electrons) and nonmetals have high electron affinities (favorable to gain electrons). Both ions achieve a stable noble gas electron configuration.

Check: Electron Transfer

Lattice Energy

Once ions form, they don't exist as isolated pairs. They arrange into a crystal lattice — a repeating 3D structure of alternating cations and anions.

What Is Lattice Energy?

Lattice energy is the energy released when gaseous ions come together to form one mole of an ionic solid:

$\text{M}^{n+}(g) + \text{X}^{m-}(g) \rightarrow \text{MX}(s) \quad \Delta H_{lattice} < 0$

A larger (more negative) lattice energy means a more stable ionic compound.

Coulomb's Law for Ions

The lattice energy is proportional to the Coulombic attraction between ions:

$E \propto \frac{q_+ \times q_-}{r_+ + r_-}$

where:

• $q_+$ and $q_-$ are the charges on the cation and anion
• $r_+$ and $r_-$ are the ionic radii

Trends in Lattice Energy

Examples

Lattice Energy Ranking

Properties of Ionic Compounds

The strong electrostatic forces in ionic crystal lattices give ionic compounds distinctive physical properties:

Why Brittle?

When a force shifts one layer of the lattice, cations end up next to cations and anions next to anions. The resulting repulsion causes the crystal to crack along that plane.

Properties of Ionic Compounds — Fill in the Blanks

Part 1 Practice — Ion Formation

Determine the charge on each ion formed and the formula of the resulting ionic compound.

1. Calcium ($Z = 20$, Group 2) loses electrons to form a cation. What is the charge on the calcium ion? (Enter as a number, e.g., 2)

2. Fluorine ($Z = 9$, Group 17) gains electrons to form an anion. What is the charge on the fluoride ion? (Enter as a number, e.g., -1)

3. What is the ratio of Ca ions to F ions in calcium fluoride? (Enter as a ratio like 1:2)

Part 2: Covalent Bonds

Part 2 of 7 — Electron Sharing

While ionic bonds involve the transfer of electrons, covalent bonds form when atoms share electrons. Covalent bonding is the dominant bond type in molecular compounds — the molecules that make up most of organic chemistry, biochemistry, and atmospheric chemistry.

Covalent Bond Formation: Electron Sharing

A covalent bond forms when two nonmetal atoms share one or more pairs of valence electrons so that each atom achieves a more stable electron configuration.

Why Share?

Nonmetals have high ionization energies — it's too costly to remove electrons completely. Instead, atoms achieve stability by sharing electrons so that each atom has access to a full valence shell (usually an octet, or a duet for hydrogen).

Example: $\text{H}_2$

Each hydrogen atom has 1 electron and needs 2 (a duet). By sharing their electrons, both atoms have access to 2 electrons:

$\text{H}\cdot + \cdot\text{H} \rightarrow \text{H—H}$

The shared pair of electrons is attracted to both nuclei simultaneously. This mutual attraction is what holds the molecule together.

The Bonding Pair

• The shared electrons spend most of their time between the two nuclei
• Both nuclei are attracted to this shared electron density
• This creates a net attractive force that holds the atoms together

Check: Covalent Bond Basics

Single, Double, and Triple Bonds

Atoms can share more than one pair of electrons:

Examples

Oxygen ($\text{O}_2$): Each oxygen has 6 valence electrons and needs 2 more for an octet. They share 2 pairs → double bond.

$\text{O=O}$

Nitrogen ($\text{N}_2$): Each nitrogen has 5 valence electrons and needs 3 more. They share 3 pairs → triple bond.

$\text{N≡N}$

Key Pattern

The number of bonds an atom typically forms = 8 minus the number of valence electrons (for atoms that follow the octet rule).

Identify the Bond Type

Bond Energy and Bond Length

Bond Energy (Bond Dissociation Energy)

Bond energy is the energy required to break one mole of a particular bond in the gas phase:

$\text{A—B}(g) \rightarrow \text{A}(g) + \text{B}(g) \quad \Delta H = \text{bond energy} > 0$

Breaking bonds always requires energy (endothermic). Forming bonds always releases energy (exothermic).

Bond Length

Bond length is the distance between the nuclei of two bonded atoms, measured at the point of minimum potential energy.

Trends: Single vs. Double vs. Triple

Why?

More shared electron pairs means:

• More electron density between the nuclei → stronger attraction → higher bond energy
• Nuclei pulled closer together → shorter bond length

Typical C–C Bond Data

Bond Energy and Length Quiz

Part 2 Practice — Bond Analysis

1. How many shared electron pairs are in a double bond?

2. A C—C single bond is 154 pm long and a C≡C triple bond is 120 pm. Which is shorter: the single or triple bond? (Enter "single" or "triple")

3. Carbon has 4 valence electrons. Using the rule (8 minus valence electrons), how many bonds does carbon typically form?

Part 3: Metallic Bonding

Part 3 of 7 — The Sea of Electrons

We've covered ionic bonds (electron transfer) and covalent bonds (electron sharing). Now we turn to the third major bond type: metallic bonds. This model explains the unique properties of metals — conductivity, malleability, luster, and more.

The Electron Sea Model

In a metallic solid, atoms are packed closely together in a regular arrangement. Unlike ionic or covalent compounds, the valence electrons in a metal are not transferred or shared between specific atoms. Instead, they are delocalized — free to move throughout the entire metal structure.

The Model

• Metal cations form a fixed lattice (the "positive cores")
• Valence electrons are delocalized, forming an "electron sea" that flows around and between the cations
• Every cation is attracted to the surrounding sea of electrons, and every electron is attracted to many cations simultaneously

What Holds It Together?

The metallic bond is the electrostatic attraction between the positively charged metal cations and the delocalized valence electrons. This is a non-directional interaction — unlike covalent bonds, which are localized between specific atom pairs.

Metallic Bond Strength

Metallic bond strength depends on:

Example Comparison

Metals with more valence electrons and smaller ionic radii tend to have stronger metallic bonds and higher melting points.

Check: Electron Sea Model

Electrical and Thermal Conductivity

Electrical Conductivity

Metals are excellent electrical conductors because their delocalized electrons can flow freely in response to an applied electric field (voltage).

• Apply a voltage → electrons drift toward the positive terminal
• The "electron sea" provides a continuous pathway for charge flow
• This is why metals are used for wires, circuits, and electrodes

Key difference from ionic compounds:

• Ionic solids do NOT conduct (ions fixed in lattice)
• Ionic liquids/solutions DO conduct (ions mobile)
• Metals conduct in both solid and liquid states

Thermal Conductivity

Metals also conduct heat well because:

1. Free electrons can carry kinetic energy rapidly through the metal
2. Vibrations in the lattice (phonons) also transfer energy, but electron transport is faster

This is why a metal spoon in hot soup gets hot quickly, while a wooden spoon stays cool.

Malleability and Ductility

Definitions

• Malleable: Can be hammered or pressed into thin sheets (e.g., aluminum foil)
• Ductile: Can be drawn into thin wires (e.g., copper wire)

Why Are Metals Malleable?

When a metal is struck:

1. Layers of cations slide past each other
2. The delocalized electron sea adjusts and continues to hold the cations together
3. The metallic bond is maintained — no fracturing

Contrast with Ionic Compounds

When an ionic crystal is struck:

1. Layers shift → like charges (cation next to cation) line up
2. Electrostatic repulsion causes the crystal to shatter
3. Ionic compounds are brittle, not malleable

Metallic Luster

Metals are shiny because the delocalized electrons can absorb and re-emit photons of light across a wide range of wavelengths. The "electron sea" interacts with incoming light, reflecting it back — producing the characteristic metallic luster.

Metallic Properties — Fill in the Blanks

Alloys

An alloy is a mixture of a metal with one or more other elements (usually metals). Alloys modify the properties of pure metals.

Types of Alloys

Why Use Alloys?

• Harder and stronger than pure metals (different-sized atoms disrupt layer sliding)
• Customizable properties — adjust composition to tune hardness, corrosion resistance, etc.
• Still maintain metallic properties (conductivity, luster)

Why Are Alloys Harder?

In a pure metal, uniform layers slide easily. In an alloy, atoms of different sizes disrupt the regular arrangement, making it harder for layers to slide past each other.

Part 3 Exit Quiz

Part 4: Bond Polarity

Part 4 of 7 — Electronegativity and Polar Bonds

Not all covalent bonds are created equal. When two atoms share electrons, the sharing may not be equal — one atom may attract the shared electrons more strongly. This creates a polar covalent bond, a concept that bridges the gap between pure covalent and ionic bonding.

Electronegativity

Electronegativity ($\chi$) is a measure of an atom's ability to attract shared electrons toward itself in a covalent bond.

The Pauling Scale

Linus Pauling developed the most widely used electronegativity scale:

Periodic Trends

• Across a period (left → right): Electronegativity increases (higher $Z_{eff}$, stronger pull on shared electrons)
• Down a group (top → bottom): Electronegativity decreases (valence electrons farther from nucleus, weaker pull)

Fluorine is the most electronegative element. Cesium and francium are the least electronegative.

Note: Noble gases are generally not assigned electronegativity values because they rarely form bonds.

Electronegativity Check

Electronegativity Difference and Bond Type

The electronegativity difference ($\Delta\chi$) between two bonded atoms determines the bond type:

$\Delta\chi = |\chi_A - \chi_B|$

⚠️ AP Note: These boundaries are approximate guidelines, not rigid cutoffs. The transition from polar covalent to ionic is a continuum, not a sharp divide. On the AP exam, focus on the concept of a spectrum rather than memorizing exact numbers.

Examples

Polar Covalent Bonds in Detail

In a polar covalent bond, electrons are shared unequally. The more electronegative atom pulls the shared electrons closer to itself.

Partial Charges

This unequal sharing creates partial charges (denoted $\delta$):

• The more electronegative atom gains a partial negative charge: $\delta^-$
• The less electronegative atom gains a partial positive charge: $\delta^+$

Example: HCl

$\overset{\delta^+}{\text{H}} — \overset{\delta^-}{\text{Cl}}$

Chlorine ($\chi = 3.0$) is more electronegative than hydrogen ($\chi = 2.1$), so the shared electrons spend more time near Cl, giving it a $\delta^-$ charge.

Dipole Moment

A polar bond has a dipole moment ($\mu$), which is a vector pointing from the positive end toward the negative end:

$\mu = q \times d$

where $q$ is the magnitude of the partial charge and $d$ is the bond length.

The larger the electronegativity difference and the longer the bond, the larger the dipole moment.

Classify These Bonds

Use electronegativity differences to classify each bond.

The Bonding Continuum

Bond types are not distinct categories — they form a continuum:

$\text{Nonpolar Covalent} \longleftrightarrow \text{Polar Covalent} \longleftrightarrow \text{Ionic}$

Percent Ionic Character

Every bond (except between identical atoms) has some degree of ionic character. As $\Delta\chi$ increases:

• The bond becomes more polar
• The electron distribution becomes more asymmetric
• Eventually, the bond is essentially ionic (electrons effectively transferred)

Important AP Concept

Even bonds classified as "ionic" have some covalent character, and vice versa. The AP exam often tests whether students understand that bonding is a spectrum, not a set of discrete categories.

Quick Rule of Thumb

• Same element → nonpolar covalent ($\text{H}_2$, $\text{O}_2$, $\text{N}_2$)
• Two different nonmetals → usually polar covalent ($\text{HCl}$, $\text{H}_2\text{O}$)
• Metal + nonmetal → usually ionic ($\text{NaCl}$, $\text{MgO}$)

Part 4 Practice — Electronegativity and Polarity

1. Calculate $\Delta\chi$ for an O—H bond. (O: $\chi = 3.5$, H: $\chi = 2.1$. Enter to 3 significant figures.)

2. In the bond H—F, which atom carries the partial negative charge ($\delta^-$)? (Enter "H" or "F")

3. Calculate $\Delta\chi$ for a Na—Cl bond. (Na: $\chi = 0.9$, Cl: $\chi = 3.0$. Enter to 3 significant figures.)

Part 5: Comparing Bond Types

Part 5 of 7 — Ionic, Covalent, and Metallic Properties Side by Side

Now that you've studied all three bond types individually, it's time to bring them together. Being able to compare and predict bond types and their resulting properties is essential for the AP Chemistry exam.

The Big Comparison Table

Key Distinctions to Remember

1. Ionic solids don't conduct; ionic liquids/solutions do
2. Metals conduct in all states (solid and liquid)
3. Molecular (covalent) compounds generally don't conduct in any state
4. Ionic and metallic compounds have high melting points; molecular compounds have low melting points

Comparing Bond Types Quiz

Predicting Bond Type from Elements

You can predict the most likely bond type based on the types of elements involved:

Decision Tree

1. Metal + nonmetal → Ionic bond

   • Large $\Delta\chi$; electron transfer
   • Examples: NaCl, CaO, KBr

2. Nonmetal + nonmetal → Covalent bond

   • Similar electronegativities; electron sharing
   • Examples: H₂O, CO₂, NH₃

3. Metal + metal → Metallic bond

   • Delocalized valence electrons
   • Examples: Fe, Cu, brass (Cu + Zn alloy)

Exceptions and Special Cases

The Polyatomic Ion Exception

Consider ammonium chloride ($\text{NH}_4\text{Cl}$):

• Inside $\text{NH}_4^+$: nitrogen and hydrogen share electrons (covalent bonds)
• Between $\text{NH}_4^+$ and $\text{Cl}^-$: electrostatic attraction (ionic bond)

Many real compounds contain both ionic and covalent bonds!

Predict the Bond Type

For each pair of elements, predict the primary bond type.

Special Case: Covalent Network Solids

Not all covalent compounds are soft with low melting points. Covalent network solids are an important exception:

What Are They?

In a covalent network solid, atoms are connected by continuous covalent bonds throughout the entire structure — there are no individual molecules.

Examples

Properties of Network Covalent Solids

• Extremely high melting points (must break strong covalent bonds)
• Very hard (diamond is the hardest natural substance)
• Do not conduct electricity (no mobile charges)
  • Exception: Graphite — delocalized electrons in its layered structure allow conduction
• Insoluble in virtually all solvents

Substance Identification Quiz

Part 5 Practice

1. Do ionic solids conduct electricity? (Enter "yes" or "no")

2. What type of bonding is present in a sample of pure copper? (Enter "ionic", "covalent", or "metallic")

3. A substance is hard, brittle, has a high melting point, and conducts when dissolved. What bond type is present? (Enter "ionic", "covalent", or "metallic")

Part 6: Problem-Solving Workshop

Part 6 of 7 — Mixed Practice with Chemical Bonds

This part is a hands-on workshop. You'll apply everything from Parts 1–5 to identify, compare, and analyze chemical bonds. Work through each section carefully — these are the types of questions you'll see on the AP exam.

Round 1: Identify the Bond Type

Classify the primary bond type in each substance.

Round 2: Match Properties to Bond Type

A substance is described. Identify its bond type.

Round 3: Lattice Energy Comparisons

Round 4: Bond Polarity Calculations

Calculate the electronegativity difference and classify each bond.

1. $\Delta\chi$ for C—Cl (C: 2.5, Cl: 3.0). Enter the value.

2. Classify the C—Cl bond as "nonpolar covalent", "polar covalent", or "ionic".

3. $\Delta\chi$ for Na—F (Na: 0.9, F: 4.0). Enter the value.

Round all answers to 3 significant figures.

Round 5: Mixed Concept Questions

Round 6: Quick Classification Drill

Part 7: Synthesis & AP Review

Part 7 of 7 — AP-Style Questions and Common Misconceptions

This final part brings everything together with AP-exam-style questions and addresses the most common misconceptions students have about chemical bonding. Mastering these will help you earn maximum points on the AP Chemistry exam.

Common Misconceptions About Bonding

❌ Misconception 1: "Ionic bonds are stronger than covalent bonds"

Reality: This is an oversimplification. Individual covalent bonds can be extremely strong (e.g., C—C in diamond, N≡N at 945 kJ/mol). The distinction is between:

• Ionic compounds: strong interactions throughout the crystal lattice → high MP
• Molecular covalent compounds: strong bonds within molecules, but weak forces between molecules → low MP

It's the intermolecular forces (not the covalent bonds) that determine melting point in molecular substances.

❌ Misconception 2: "Electrons are completely transferred in ionic bonds"

Reality: Even ionic bonds have some covalent character. The electron cloud of the anion is distorted toward the cation (called polarization). Bonding is a continuum, not two separate boxes.

❌ Misconception 3: "Ionic compounds exist as discrete molecules"

Reality: Ionic compounds form extended crystal lattices, not individual molecules. The formula $\text{NaCl}$ represents the simplest ratio of ions, not a molecule. We call it a formula unit.

❌ Misconception 4: "All covalent compounds have low melting points"

Reality: Molecular covalent compounds generally have low MPs, but covalent network solids (diamond, SiO₂) have extremely high melting points because you must break covalent bonds throughout the structure.

❌ Misconception 5: "Metals conduct because they have ionic bonds"

Reality: Metals conduct because of delocalized electrons in the electron sea. They have metallic bonds, not ionic bonds. Ionic solids do NOT conduct electricity.

Misconception Buster Quiz

AP-Style Free Response Strategies

On the AP Chemistry exam, bonding questions often appear in the free-response section. Here are key strategies:

What They Test

1. Identifying bond type from a compound formula or element combination
2. Explaining properties (MP, conductivity, hardness) in terms of bonding
3. Comparing substances using Coulomb's law or bond polarity
4. Lattice energy ranking using ion charge and size

How to Earn Full Points

Always connect your answer to the underlying model:

The Coulomb's Law Argument

When comparing ionic compounds, always invoke Coulomb's law:

$F \propto \frac{q_1 \times q_2}{r^2}$

Specify which charges and which radii you're comparing. The AP exam rewards specific, quantitative reasoning.

AP-Style Multiple Choice — Set 1

AP-Style Multiple Choice — Set 2

Comprehensive Classification

Identify the primary bond type and predict a key property.

Final Challenge — Apply What You've Learned

1. What type of bond forms between two fluorine atoms in $\text{F}_2$? (Enter "nonpolar covalent", "polar covalent", or "ionic")

2. In the compound $\text{LiF}$, which element gains the electron? (Enter "Li" or "F")

3. How many valence electrons does nitrogen ($Z = 7$) have?